Chemistry
Redox Potential Of Transition Metals
The redox potential of transition metals refers to their ability to undergo oxidation-reduction reactions. It is a measure of the tendency of a transition metal to gain or lose electrons. Transition metals with higher redox potentials are more likely to be oxidized, while those with lower redox potentials are more likely to be reduced. This property is important in various chemical and industrial processes.
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8 Key excerpts on "Redox Potential Of Transition Metals"
eBook - PDF- Rozina Khattak(Author)
- 2020(Publication Date)
- IntechOpen(Publisher)
The amount of literature on this topic is very large and is beyond the scope of this manuscript to discuss it in any detail. As pointed out above, there are several reviews on this subject [6, 18, 23, 29 – 31], so discussion will be focused on using the complex redox potential as a reactivity descriptor. Randin was the first to attempt to rationalize the catalytic activity of MN4 complexes [32] in 1974 later followed by Beck in 1977 [33] using simply the redox potential of the MN4 complex as a reactivity indicator. A first theoretical explana-tion based on these terms was provided by Ulstrup in 1977 [34]. According to Randin and Beck, during the binding of O 2 to the metal ion in the complex, the metal is partially oxidized, thereby reducing the O 2 molecule according to the: M II ð Þ þ O 2 ! M III ð Þ O � 2 � � (14) M III ð Þ O � 2 � � ! products þ M III ð Þ (15) In order to account for supplementary experimental evidence, a somewhat modified model was proposed by Beck in which the central metal ion might also be only partially oxidized. According to reactions in Eqs. (14) and (15), the potential at Figure 10. Mathematical simulation of the variation of fractional coverage θ Fe(II) on the electrode surface as a function of potential for two hypothetical values of E Fe(II)/(I) and E Fe(III)/(II) . Simulation involved the Nernst equation applied to surface confined species, assuming ideal behaviour (adapted from [26]). 61 Redox Potentials as Reactivity Descriptors in Electrochemistry DOI: http://dx.doi.org/10.5772/intechopen.89883 which O 2 is reduced should be closely related to the M(III)/M(II) redox potential of the central metal ion.
- Robert Hanzlik(Author)
- 2012(Publication Date)
- Academic Press(Publisher)
With the exception of lithium there is a general trend in the redox potential exhibited by nontransition metals: The redox potential (emf) increases toward the left of a row and toward the bottom of a family in the periodic table. Among the transition metals, trends in redox potentials are complicated by FACTORS INFLUENCING THE REDOX POTENTIAL OF METALS 141 crystal field stabilization energy effects on hydration energy and the problems associated with the changing of d orbital energy levels with atomic number. Among a vertical triad, however, the heavier atoms usually are more stable in higher oxidation states, and the lighter atoms, in their lower oxidation states. Compare, for example, Cr(III) vs W0 4 2 , Fe(III) vs Os0 4 , and Ni(II) vs Pt(IV). LIGAND EFFECTS The lower the energy of an orbital, the more stable an electron in that orbital will be. Because of the way in which ligands split the energies^of the d orbitals, the type of orbital an electron enters or leaves during redox will have an effect on the redox potential. This can be seen by considering the Fe(II)/Fe(III) couple for the complexes shown in Diagram 7-1. 3d configuration Fe(III) + e~ = Fe(II) ++ +++ ++ low spin Fe complex E 0 Fe(ö-phen) 3 2+ ' 3+ +1.14 Fe(CN) 6 4 -' 3 +0.36 +++ +++ high spin Fe(H 2 0) 6 2 Diagram 7-1 +0.77 Comparison of the o-phenanthroline (o-phen) complex with the aquo com-plex illustrates the effect of electron configuration on redox potential of com-plexes with equal net ionic charge. In each case the electron enters a t 2g orbital, but the orbital is of lower energy in the o-phen complex than in the aquo com-plex because the Δ value for o-phen is much larger than the Δ value for water. Although the cyano complex has the same electronic configuration as the o-phen complex, its redox potential is far lower, mainly because it has a net charge of 3—, rather than 3+ as the others have, and this large accumulation of negative charge tends to resist the entrance of another electron.
eBook - PDF- Bernard Moody(Author)
- 2013(Publication Date)
- Arnold(Publisher)
The magnitude of tion must relate to the passage of one Faraday of this potential depends on the net tendency of the electric charge. system to absorb electrons from the platinum elec-If the equation had been written the other way trode or to donate them to it, round Fe 3+ 2+ + e ^ Fe ^-Cu + £Zn 2+ iCu 2+ + ±Zn The potential is called a redox potential. At 25°C, this would imply the diagram using concentrations of iron(n) and iron(m) salts Cu | Cu 2+ | Zn 2+ | Zn which are both 1.0 M, the standard redox potential Table 7.2 Standard electrode potentials (ctd. standard redox potentials at 298K (25°C) Pt electrode in Electrode process E 2 ts/V H + ; H 2 FT + e ^ i H 2 0.000 Sn 4+ ; Sn 2+ iSn 4+ + e~ ^ iSn 2+ +0.15 Cu 2+ ; Cu + Cu 2+ + e ^ C u + +0.153 Fe(CN) 6 3 -; Fe(CN) 6 4 Fe(CN) 6 3 -+ e ^ Fe(CN) 6 4 -+0.360 ^Fe 2 + Fe 3+ ; Fe 2+ Fe 3+ + e +0.771 2+ 2+ Hg 2+ ; Hg 2 Hg 2+ + e -iHg 2 +0.920 Cr 2 0 7 2 -; Cr 3+ iCr 2 0 7 2 + lH + + e~ ^ iCr 3+ + £H 2 0 + 1.33 ^Ce 3 + Ce 4+ ; Ce 3+ Ce 4+ + e + 1.45 Mn(V; Mn 2+ ^MnCV + |H + + e ^ ^Mn 2+ + £H 2 0 + 1.52 The physical states are implied by the standard conditions 120 Oxidation, reduction and electrochemical processes may be determined on the standard hydrogen scale. Standard redox potentials are listed in Table 7.2. They are separated from those given in Table 7.1 purely for convenience: both tables show redox potentials. The relative positions of two systems indicate the possibility of a reaction between them but give no information about the rate of reaction, the influence of catalysts, or the other variables such as complex ion or insoluble product forma-tion.
eBook - ePubInorganic Chemistry for Geochemistry and Environmental Sciences
Fundamentals and Applications
- George W. Luther, III(Authors)
- 2016(Publication Date)
- Wiley(Publisher)
(an inner sphere electron transfer process).10.2 Factors Governing Metal Speciation in the Environment and in Organisms
Many factors affect the reactivity of a metal ion in the environment. Section 1.8.2 indicated that the oxidation state of a metal [e.g., Fe(II) versus Fe(III)] depends on the redox condition of the environment. For example, once dissolved and are consumed, microbes decompose organic matter with other oxidants such as and FeOOH, resulting in reduction to and . Metal ions can exist as inorganic complexes with the following common inorganic ligands: chloride, carbonate, sulfate, phosphate, and sulfide. However, there are a variety of organic ligands that are naturally produced and that outcompete the inorganic ligands for bonding with metal ions. Ligands also affect the redox and spin state of a metal couple such as Fe(III)/Fe(II) (Section 8.7.3 for Co; Table 10.1 for Fe), and nature uses ligand–metal bonding to affect reactivity and catalysis. The left of Figure 10.1 shows potentials for several reduction couples that are important in life processes (sometimes termed a redox spectrum). An oxidized partner of a redox couple at the top such as can be reduced by the reduced partner of any couple below it as . Thermodynamically, the couple is one of the most efficient, but it also shows the energy needed for water splitting (the reverse reaction that requires photochemistry; Section 10.7 ). The FeS proteins (see Section 12.6.4 ) have different oxidation states and ligand attachments that tune their redox potential over 1 volt; reduction of to is possible for some redox centers (ferredoxins). Selected aspects of the redox chemistry for the metal redox centers in Figure 10.1 are described below and in Chapter 12
eBook - PDF- A.T. Diplock, M.C.R. Symons, C.A. Rice-Evans(Authors)
- 1991(Publication Date)
- Elsevier Science(Publisher)
CHAPTER 4 Transition metal complexes as sources of radicals 4.1. Introduction We start with a general outline of the many roles played by transition metal (TM) complexes as redox agents and as sources of radicals. This is followed by a range of examples of processes thought to be of im- portance in biological systems. These include reactions in which radi- cals are produced which can lead to biological damage, as well as sys- tems in which their formation is part of the biologically intended reaction. 4.2. Transition metal ions in biological systems Only transition-metal ions are redox-active and hence are possible sources of radicals. Ions such as Mg2+, Ca2+, Zn2+ etc. are not im- portant in this context. By far the most important is iron, with copper, molybdenum, cobalt and nickel also participating on a more minor scale. Of course, any transition metal ion may be ingested, and hence may be a fortuitous source of radical formation and damage. We focus attention on iron for illustrative purposes. The concentrations of iron as simple solvated ions, Fe2+,q or Fe3+,q, are maintained at extremely low levels because of their da- maging ability in the presence of oxygen and hydrogen peroxide. In transferrin, an important iron scavenger, the iron is well-protected and is not involved in redox chemistry. Also, in the major storage pro- teins, ferritin and haemosiderin, the iron is present in crystalline mate- rial inside the protein shell, and is well protected from reaction. 102 TECHNIQUES IN FREE RADICAL RESEARCH 4.2.1. Factors governing redox behaviour Redox potentials for a given metal ion are strongly controlled by the types of ligands involved, their orientation (particularly strongly con- trolled in proteins), the gain or loss of ligands and electron-delocaliza- tion onto ligands (Table 4.1).
eBook - PDF- F Dwyer(Author)
- 2012(Publication Date)
- Academic Press(Publisher)
A simple example is provided by the [Ru(bipy) 2 -C l 2 ] ° /+ couple, the reduced state of which is very sparingly soluble in water. Iodine, ferric ion, and bromine fail to effect oxidation, but chlorine oxidizes the substance. The potential is therefore between —0.9 and —1.35 volts, allowing for about 0.2 volt for complete oxidation. If bromide ion is now added to the oxidized form of the couple, it is found that bromine is not produced, hence the potential must be slightly more positive than the B r _ / B r 2 couple, ( — 1.02 volts), i.e., approximately —0.95 volt. These semi-quantitative estimates are only valid with reversible systems; failure to effect oxidation or reduction with a particular reagent may be due to mechanistic reasons, and not to the (presumed) unfavorable free-energy change. The redox potentials of complex couples which dissociate reversibly can be calculated from their stability constants if E° for the aqueous metal ion couples is known (Perrin, 1959; Hawkins and Perrin, 1962). The theoretical basis for this will be evident if the following system is con-sidered at equilibrium : Equation (4) gives the redox potential E of the complex couple MLJ / M L x ( n + 1 ) + . At equilibrium, when E = 0, this becomes B . R E D O X POTENTIALS FROM STABILITY CONSTANTS fix M ( n + D + + xh M L X < » + 1 > + +eJT . t l + e E° = or E° = ^± i n 6. FUNCTIONS OF DONOR ATOM AND LIGAND 2 4 5 Now QIT = [ a MLs(+D+] [a M («+i)+][a L ] x and R , [«ML, + ] Px [a M n + ][a L ]-hence «ML,«+ Px #M« + and T aM«+ R Px At equilibrium also RT ^ M«+/M(« + D+ = ~FT In /< aM«+ whence i ? 0 (complex couple) = E° (hydrated metal ion couple) H—^- In r Px Thus the thermodynamic E° for a complex couple can be obtained from the stability constants of the complexes extrapolated to infinite dilution pro-vided E° for the aqueous metal ions is known.
eBook - ePubMetal Ions in Biological Systems
Volume 36: Interrelations Between Free Radicals and Metal Ions in Life Processes
- Astrid Sigel, Helmut Sigel(Authors)
- 2018(Publication Date)
- CRC Press(Publisher)
Free radicals and transition metal ions are ubiquitous in biological systems. The variable valency of transition metal ions plays a key role in the catalysis of redox processes in which the oxidation states of the metal are inherently unstable and/or have only a transient existence during the reaction sequence. In many cases where the overall reaction involves multiple electron transfer, the detailed mechanism frequently comprises a sequence of one-electron steps. Free radicals, being atoms or molecules with an unpaired electron, are the archetypal one-electron redox reagents, so that the elementary step in the interaction of free radicals with transition metal ions is a one-electron process. The elucidation of the complex mechanisms that take place in nature can be advanced to a great extent if one is able to measure in isolation these individual steps. To achieve this goal use is made of the fast time-resolved methods of pulse radiolysis and flash photolysis to generate the desired free radicals and to measure their rates of reaction with the species of interest. Often these species are model compounds rather than the biological systems themselves.Previous surveys of uncommon oxidation states of metal ions can be found in the literature [1 , 2 , 3 ].1.2. Interaction of Free Radicals with Transition Metal IonsA characteristic property of the transition metals in solution is that they are coordinated to ligands, so that one must consider the complete entity—metal center and ligands —in its reaction with free radicals. The possible reactions are of two main types: outer-sphere electron transfer, in which an electron is transferred between the free radical and the metal center with the ligands of the latter remaining intact; and inner-sphere electron transfer, whereby the ligands are involved. In the latter case the reaction may occur first at a ligand, turning it into a radical, followed by electron transfer to or from the metal center, or the free radical may itself become bonded to the metal ion, either by exchanging with one of the ligands or by increasing the coordination number of the metal ion by one, before electron transfer takes place. In some cases the free radical remains bound to the metal center and the oxidation number of the metal is not well defined. Examples of each of these kinds of reaction are described in Secs. 5 and 6.
- Henry Freiser, Monika Freiser(Authors)
- 1992(Publication Date)
- CRC Press(Publisher)
Chapter 7 Oxidation Reduction Equi l i bria Oxidation-reduction (redox) reactions involve electron transfer and therefore can be used to produce electrical work. This is accomplished by suitably separating the reaction components into two halfcells linked in a way that electron transfer must occur through an external circuit. Such a system is called a galvanic cell. This link between electrical and chemical transformations, constituting the field of electrochemistry, is vital to understanding and controlling many phenomena of interest to a wide range of scientists. It provides analytica l chemists with powerful methods of monitoring chemical reactions by rapid, precise, and convenient electrochem-ical measurements of concentrations over a wide range. Transfer of chemical to electrical energy in batteries is responsible for an important form of clean energy. Electroplating represents another application of wide interest. Electrochemistry has even been exploited as an apocryphal means of accomplishing cold fusion. The electromotive force (emf) of the cell is a measure, in volts, of the driving force of the chemical reaction involved. Although one cannot measure directly the voltage of each of the two halfcells, the emf of the cell is the voltage difference between the two halfcells. As will be described in detail below, halfcell potentials can be quantitatively evaluated by adopting one halfcell as a standard and measuring all the others by the potential of the cell, Ex -Eref. The equilibrium constant of the chemical reaction is another measure of this driving force. The electromotive force of the cell and the equilibrium constant of the chemical reaction involved are therefore related as will be shown below. As a matter of common practice, however, redox equilibria are more often characterized in terms of electromotive force values rather than equilibrium constants.
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